Chemical Bonding - Part I

Gap-Fill Exercise
Content © 2006 Melanie Cecere; Authors of Activity: Melanie Cecere & Tami Maloney; All rights reserved. No commercial, for-profit use of this material is allowed. E-mail comments and questions to Tami Maloney.

Fill in all the gaps, then press "Check" to check your answers. Use the "Hint" button to get a free letter if an answer is giving you trouble. You can also click on the "[?]" button to get a clue. Note that you will lose points if you ask for hints or clues!
seldom exist as independent particles in nature. A chemical is a mutual attraction between the nuclei and valence electrons of different atoms that binds the atoms together. As independent particles, atoms are at relatively potential energy. By bonding with each other, atoms in potential energy, thereby creating more stable arrangements of matter.
- When atoms bond, their electrons are redistributed. The way they are redistributed determines the type of bonding. bonding involves transfer of electrons. It results from the electrical attraction between large numbers of and . bonding results from sharing of electron pairs between two atoms.
- Bonding between atoms of different elements is never purely ionic and is rarely purely covalent. It usually falls somewhere between these two extremes. is a measure of an atom's ability to attract electrons in a bond. The degree to which bonding between atoms of two elements is ionic or covalent can be estimated by calculating the difference in the elements' electronegativities. As the difference between electronegativities becomes greater, bonds become increasingly . When the difference is very great, bonds are predominantly .

A covalent bond is a covalent bond in which the bonding electrons are equally by the bonded atoms, resulting in a balanced distribution of electrical charge. In bonds with significantly different electronegativities, the electrons are more strongly attracted by the more atom. Such bonds are covalent (those with an uneven distribution of charge). A polar-covalent bond is a covalent bond in which the bonded atoms have an unequal attraction for the shared electrons. In HCl, the electronegativity difference between chlorine and hydrogen is 3.0 - 2.1 = 0.9, indicating a polar-covalent bond. The electrons are closer to the more electronegative atom than to the hydrogen atom. Consequently, the chlorine end of the bond has a partial charge, and the hydrogen end of the bond has an equal partial charge.

A is a neutral group of atoms that are held together by covalent bonds. It is a single unit capable of existing on its own. A molecular compound is a chemical compound whose simplest units are . The chemical indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts. A formula shows the types and numbers of atoms combined in a single molecule of a molecular compound. For example, H2O is the molecular formula for , indicating that there is one atom joined by separate covalent bonds to two atoms. A molecule is one that contains only two atoms.

Nature favors chemical bonding because most atoms are at potential energy when bonded to other atoms than they are as independent particles. As atoms near each other, their charged particles begin to interact. The approaching nuclei and electrons are attracted to each other, which corresponds to a in the total potential energy of the atoms. At the same time, the two nuclei each other and the two electrons each other, which result in an in potential energy. The relative strength of attraction and repulsion between the charged particles depends on the distance separating the atoms. There is an optimum distance where their attractive forces equals the repulsion between the like charges and this corresponds to the bottom of the valley on the potential energy curve in a stable covalent bond.

The bond is the distance between two bonded atoms at their potential energy or the average distance between two bonded atoms. The bond is the amount of energy required to break a chemical bond and form neutral atoms. It is the same amount of energy released when the bond was formed, measured in kJ/mol. The bond length is proportional to bond strength. Bond lengths and bond energies vary with the types of atoms that have combined. Even the energy of a bond between the same two types of atoms varies somewhat, depending on what other bonds the atoms have formed.

Unlike other atoms, the atoms exist independently in nature. They possess a minimum of energy existing on their own because of the special of their electron configurations. Other main group atoms can effectively fill their outermost s and p orbitals with electrons by sharing electrons through bonding. Such bond formation follows the rule: chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level.
There are exceptions. Hydrogen forms bonds in which it is surrounded by only electrons. Boron, B, has just 3 electrons. Because electron pairs are shared in covalent bonds, boron tends to form bonds in which it is surrounded by 6 electrons. Other elements can be surrounded by more than 8 electrons when they combine with the highly elements F, Cl, and O. In cases of expanded valence, bonding involves electrons in d orbitals as well as in s and p orbitals.

- Covalent-Network Bonding: All the covalent compounds that you have read about to this point are molecular. They consist of many identical molecules bound together by forces acting between the molecules. There are many covalently bonded compounds that do not contain individual , but instead can be pictured as continuous, 3-dimensional networks of bonded atoms. Two examples are SiO2 (sand, quartz), and diamonds.

Ionic Bonding and Ionic Compounds. Most of the rocks and minerals that make up the earth's crust consist of positive and negative ions held together by bonding. A familiar example of an ionic compound is NaCl, common table salt or in nature as rock salt. A ion, Na+, has a charge of +1. A ion, Cl-, has a charge of -1. These combine in a 1 to 1 ratio (Na+Cl-) so that each charge is balanced by a negative charge. Because chemists are aware of this balance of charge, the formula is usually written simply as NaCl. An ionic is composed of positive and negative ions that are combined so that the numbers of positive and negative charges are .

- Ionic bonding. The simplest forms of ionic bonding come from a at the far left of the periodic table with a at the right. Metals lose electrons because of low energy. Nonmetals gain electrons because of a strong for electrons. in orbitals in the valence shell of the metal are transferred to orbitals in the valence shell of the nonmetal. Each gains or loses electrons to obtain the stable octet of the noble gas nearest the element. Except for , the electron configuration of the valence shell of all noble gases is s2p6. electrons rearrange themselves during a chemical reaction so that each atom has a stable octet. This is called the rule or "rule of 8".

Nature favors arrangements in which potential is minimized. In an ionic crystal, ions minimize their potential energy by combining in an orderly arrangement known as a crystal . The attractive forces at work within an ionic crystal include those between oppositely charged and those between the nuclei and electrons of adjacent ions.
- Ionic crystals. Anions and cations are packed into a regular pattern called an ionic or crystal lattice. The fundamental units of the lattice are ions. The ions are held in place by strong attractions of surrounding ions; therefore, they have very melting points. When melted or dissolved in water, the structure collapses and the ions are free. Only then can they conduct . They cannot conduct in the solid phase because they are immobile. Other examples of such ionic compounds are BaCl2, KOH, and CaCO2. To compare bond strengths in ionic compounds, chemists compare the amounts of energy when separated ions in a gas come together to form a crystalline solid. energy is the energy released when one mole of an ionic crystalline compound is formed from a gaseous ion.

The force that holds ions together in compounds is a very strong overall attraction between positive and negative charges. In a compound, the covalent bonds of the atoms making up each molecule are also strong. But the forces of attraction between molecules are much than the forces of ionic bonding. This difference in the strength of attraction between the basic units of molecular and ionic compounds gives rise to different properties in the two types of compounds.

Ionic compounds are hard, but , because even a slight shift of one row of ions relative to another causes a large buildup of repulsive forces. The strong forces make it difficult for one layer to move relative to another, causing ionic compounds to be hard.
Certain atoms bond covalently with each other to form a group of atoms that has both molecular and properties. A ion is a charged group of covalently bonded atoms. They behave as single and are very stable.

Metallic Bonding. Metals are excellent electrical conductors in the state due to mobile valence electrons of the metal atoms. The highest energy levels of most metal atoms are occupied by very few electrons. Most orbitals and some of the inner d orbitals are vacant in the metals. The vacant orbitals overlap and allow the outer electrons to roam freely throughout the entire metal. The electrons are delocalized, which means they do not belong to any one atom but move freely about the metal's network of empty . These mobile electrons form a sea of electrons around the metal atoms. bonding is the chemical bonding that results from the attraction between metal atoms and the surrounding of electrons.
- Metallic Properties. The freedom of motion of in a network of metal atoms accounts for the high electrical and thermal conductivity characteristics of metals. Because metals contain many orbitals separated by extremely small energy differences, metals can absorb a wide range of frequencies. The absorption and emission of light accounts for the appearance of metals.
- Metallic Bonding. Metals have only 1 or 2 electrons and ionization energies. The valence electrons are not tightly bonded and belong to the whole metal crystal. The electrons are shared by all of the atoms of the metal. Metals can be thought of as positive ions immersed in an "atmosphere" or "sea" of mobile electrons. The mobile electrons exert an attractive force on the positive ions to fix their positions-the attractive force that binds the metal atoms is called a metallic bond.
- Properties of Metals: Metals are (can be hammered into a shape or a thin foil); (can be stretched into a thin wire); (can be cut into sections); good conductors of heat and electricity (because of mobile ); and have high luster (as a result of the uniform way valence electrons absorb and re-emit ). Hard metals have a strong binding force; soft metals have a weak binding force. Malleability, ductility, and sectility all result from the uniform attraction between the electrons and the ions-the ions can change position without breaking up the essential structure.